How do acids and alkalis react?

How do acids and alkalis react?

Demonstration

This activity links the neutralisation of an acid by an alkali to the changes in ionic concentrations that result from the reaction between hydrogen ions and hydroxide ions.
 

Lesson organisation

The practical involves a mildly toxic alkaline solution, barium hydroxide. It needs careful preparation and manipulation, involving a conductimetric titration. A clear demonstration is more likely to lead to successful learning than a class activity that will need a lot of teacher support.

A good demonstration should take about 25 mins.

Chemicals

Each demonstration requires:

Barium hydroxide solution, 0.10 M, (HARMFUL, IRRITANT), about 200 cm³ 

Dilute sulfuric acid, 1.0 M, (IRRITANT), about 25 cm³

Purified (distilled or deionised) water

Phenolphthalein indicator solution (HIGHLY FLAMMABLE)

Apparatus

Eye protection for the teacher and for any members of class who assist at the bench

Test-tube (150 x 25 mm)

Test-tube rack

Beakers (100 cm³), 2

Measuring cylinder (100 cm³)

Glass rod

Burette (50 cm³)

Clamps (2) and stand

Small funnel (for filling the burette)

White tile (for standing beaker on during titration) and white card background (for class visibility)

Pair of carbon electrodes (Note 1) in holder, with 4 mm plug adapters

Plug leads (4 mm plug at each end), 4

Bulb (12 V) in holder

AC demonstration ammeter

Low voltage AC supply (Note 2) 

Health & Safety and Technical notes

Wear eye protection. 

Barium hydroxide solution, Ba(OH)₂(aq), (HARMFUL, IRRITANT at concentration used) - see CLEAPSS Hazcard. Solid barium hydroxide (CORROSIVE) contains water of crystallisation, and reacts with carbon dioxide from the air while in storage.  It is only slightly soluble in water (maximum 4 g in 100 cm³) but the solution is much more alkaline than limewater. Make 250 cm³ of the solution. Purified water should be boiled to remove carbon dioxide before being added to solid barium hydroxide. Once prepared the solution is very sensitive to carbon dioxide and immediately goes cloudy (barium carbonate) when exposed to the atmosphere. All fresh solutions, which will be alkaline and may irritate sensitive skin, must be protected with a soda lime guard tube. For all these reasons, it is important to check the concentration of the solution before the demonstration to ensure it is reasonably close to 0.10 M – an exact concentration is not necessary. To do this titrate 50 cm³ of the solution with the dilute sulfuric acid (1.0 M), to be used in the experiment, using phenolphthalein indicator (two drops). A titre value between 4 cm³ and 6 cm³ is acceptable.

Dilute sulfuric acid, H₂SO₄(aq), (IRRITANT at concentration used) - see CLEAPSS Hazcard and CLEAPSS Recipe Book.

Phenolphthalein indicator solution (HIGHLY FLAMMABLE) - see CLEAPSS Hazcard and CLEAPSS Recipe Book. Phenolphthalein indicator solution should be provided in a dropper bottle.

1 The two carbon electrodes need to be mounted securely in a holder that keeps them parallel. If available, a 4 mm plug adapter should be fitted to the top of each electrode. If not available, crocodile clips can be used instead, but you will need a cardboard or plastic separator between the clips to avoid accidental short-circuiting.

2 The low voltage power supply should be a variable low voltage unit capable of supplying alternating current (AC) at about 12 V when connected through a 12 V bulb and the electrodes dipped in solution in series. A demonstration AC ammeter should be included if available, and if the 12 V bulb fails to light brightly enough for the class to see when tested as in Stage 1 below, the ammeter is essential.

 

Procedure

Stage 1
a Mix equal volumes of dilute sulfuric acid and barium hydroxide solution in a test-tube to observe what happens.

b Add 50 cm³ of barium hydroxide solution to one beaker, and add 2–4 drops of phenolphthalein indicator solution to show the solution is alkaline.

c Dip the electrodes into this solution to demonstrate that it conducts electricity. The AC supply should ensure there is no electrolysis.

d Rinse the electrodes with purified water. Now test the sulfuric acid in a second beaker to show that it conducts electricity.

Stage 2
e Fill the burette with 1.0 M sulfuric acid to the zero mark. Fix the burette securely over the beaker containing 50 cm³ of 0.1 M barium hydroxide solution, ready for titration.

f Clamp the electrode assembly firmly at one side of the beaker so the electrodes dip into the full depth of the solution, and connect it to the rest of the test circuit. Place the stirring rod in the solution.

g Switch on the supply, note the ammeter reading and the bulb brightness.

h Add sulfuric acid from the burette, 0.5 cm³ at a time, with stirring. After each addition, note the ammeter reading and the bulb brightness, and look for a permanent change in the indicator colour.

i At the indicator end-point note the volume of acid added, the ammeter reading and the bulb brightness.

j Continue to add portions of acid and note the ammeter reading and bulb brightness until the change on further addition is minimal.

Teaching notes

This experiment provides important evidence for the simple hydrogen ion theory of acidity, and for the ionic nature of the neutralisation reaction with hydroxide ions. Thus it forms a natural part of a sequence of experiments in which this theoretical model can be built up for students. This experiment is not likely to be useful on its own.

It also depends on students’ understanding of ionic theory in general, and their appreciation that the conductivity of a solution depends on the concentration of ions in the solution. Alternating current is used rather than direct current to avoid electrolysis taking place, at least to any extent that would affect the outcome of the experiment.

Although this is likely to be a demonstration for most students, some teachers may wish to use it as a class experiment, possibly with older students. Safety issues are relatively minor, essentially using dilute sulfuric acid (1.0 M) and barium hydroxide solution (0.1 M). The latter should be treated as HARMFUL for students, even at this low concentration.

The reason for using such different concentrations in a conductimetric titration is to minimise the decrease in conductivity caused by increasing the volume of water present as the titration proceeds, apart from the changes in the total number of ions present.

Here are answers to the questions
1 This revises the students’ understanding of ions, and their ability to identify the ions that are present (barium cations and hydroxide anions in barium hydroxide, and hydrogen cations and sulfate anions in sulfuric acid). The symbols for these ions are:

Ba²⁺, OH^-, H⁺, SO₄^2-

2 The ions are the current carriers in solution, and conductivity depends on the concentration of ions. In asking which ions are removed in the formation of barium sulfate, there is an opportunity to write the ionic equation for the precipitation reaction

Ba²⁺(aq) + SO₄^2-(aq) → BaSO₄(s)

and hence to identify a fall in the total ion concentration which is reflected in a fall in conductivity.

3 Having dealt with the barium and sulfate ions, the students are then in a position to focus on what else is happening, starting with identifying hydrogen cations and hydroxide anions as potential reactants. These ions do react and the reaction is simple.

4 The ionic equation for the reaction between these ions is:

H⁺(aq) + OH^-(aq) → H₂O(l)

5 Any reactant added once an end-point has been reached is ‘surplus’, and so ions are again present and the conduction of the solution rises again.

It is unlikely that the end-point will be marked by zero conductivity. The titration method used is too crude to find the exact point at which almost no ions are present, and even a drop of sulfuric acid added nearly at the end-point will take the reaction past the end-point.

How can hardness in water be removed?

How can hardness in water be removed?

Demonstration and Class practical

Temporarily hard water is made by bubbling carbon dioxide through limewater for some time. Temporarily and permanently hard water are boiled, have sodium carbonate added and are subjected to ion exchange. The hardness of the solutions, before and after, is tested using soap solution.
 

Lesson organisation

This experiment is designed as a combination of demonstration and class practical work. The making and boiling of temporarily hard water is best done on a larger scale than in a test-tube, and ion exchange columns are tedious to set up in quantity. Also, the students already have a bewildering array of colourless solutions with which to deal. Adding more might cause some to get very confused. With small groups, however, the preparation and boiling of temporarily hard water and the ion exchange could be done by the students.

If the suggested method here is used, the beakers of solutions should be labelled A to F and each should have a dropping pipette. Students should bring up their test-tube racks and move along the solutions, placing 1 cm depth of each solution in the corresponding test-tube.

Do not let them mix the dropping pipettes nor move the stock bottles around.

The demonstration plus the student practical will take about one hour.

Chemicals

The teacher will require:

Calcium sulfate solution, 300 cm³ (Note 4)

Limewater (IRRITANT), 150 cm³

Sodium zeolite, about 5 g (Note 2),

Marble chips

Dilute hydrochloric acid, 2 M (IRRITANT)

Each student or group of students will need:

Solutions labelled as shown below, 100 cm³ each

A Temporarily hard water*

B Permanently hard water (Note 4)

C Temporarily hard water that has been boiled and filtered*

D Temporarily hard water that has passed through an ion exchange column*

E Permanently hard water that has passed through an ion exchange column*

F Distilled water (or deionised water)

* These are prepared by the teacher in the demonstration part of the experiment

Soap solution in IDA (Industrially Denatured Alcohol) (HIGHLY FLAMMABLE, HARMFUL),
10 cm³ (Note 5)

Sodium carbonate-10-water (IRRITANT), about 1 g

Apparatus

The teacher will require:

Eye protection

Beakers (500 cm³), 3

Bunsen burner

Tripod and gauze

Heat resistant mat

Vacuum filtration apparatus (Note 1)

Ion exchange apparatus (Note 2)

Gas generator for carbon dioxide (Note 3), or a CO₂ cylinder

Each student or group of students will need:

Test-tubes, 9

Test-tube rack

Labels, for test tubes

Beaker (100 cm³)

Dropping pipette

Spatula

Health & Safety and Technical notes

Wear eye protection throughout. 

Calcium sulfate, CaSO₄.2H₂O(s) - see CLEAPSS Hazcard.

Limewater (calcium hydroxide solution), Ca(OH)₂(aq), (treat as IRRITANT) - see CLEAPSS Hazcard and CLEAPSS Recipe Book.

Sodium zeolite 

Sodium carbonate-10-water, Na₂CO₃.10H₂O(s), (IRRITANT) - see CLEAPSS Hazcard. 

Marble chips (calcium carbonate), CaCO3(s) - see CLEAPSS Hazcard.

Hydrochloric acid, (IRRITANT) - see CLEAPSS Hazcard and CLEAPSS Recipe Book.  

1 Vacuum filtration apparatus. See CLEAPSS Laboratory Handbook.

2 Ion exchange apparatus. The sodium zeolite should be soaked in deionised water for 24 hours before use. (Dry resin would expand and crack the tube). A cotton wool plug should then be placed at the bottom of a tube with a tap* and the resin added (as a slurry) above the cotton wool. The resin must be kept covered in deionised water until the column is required. Some hard water is poured into the tube above the deionised water and the tap is opened. More hard water is added as the softened water is collected in a beaker below the tap.

* a burette will do, provided the tap is removable, allowing the cotton wool plug to be pushed out with a rod.

3 Gas generator for carbon dioxide. See Standard Techniques:

4 Calcium sulfate solution. Stir a spatula or two of calcium sulfate dihydrate into distilled water until no more will dissolve (it is not very soluble). Allow to stand and decant off the clear, saturated solution. Dilute it with an equal volume of distilled or deionised water to make the stock solution of permanently hard water.

5 Soap solution in ‘ethanol’ (Industrial Denatured Alcohol, IDA – see CLEAPSS Hazcard (HIGHLY FLAMMABLE, HARMFUL) can be purchased or made up – see CLEAPSS Recipe Book.

Procedure

Demonstration
a Dilute about 150 cm³ of limewater with an equal volume of distilled or deionised water. Pass in carbon dioxide, taking care that the gas carries over no acid spray (from the reaction between the marble chips and the acid). A milky precipitate of calcium carbonate soon forms. Continue the passage of gas until all the precipitate dissolves, giving a solution of calcium hydrogencarbonate. This is temporarily hard water.

b Place about half of the temporarily hard water in a beaker and boil it for about 5 minutes. Filter, using vacuum filtration apparatus.

c Scrape some of the solid residue from b into a test-tube and add dilute hydrochloric acid. Fizzing should show that the solid is a carbonate (calcium carbonate).

d Boil about the same quantity of permanently hard water (to that used in b) in another beaker. Show that there is no precipitate. Allow the solution to cool until it is safe to handle.

e Set up two ion exchange columns containing sodium zeolite (see note 3). Pour about half of the temporarily hard water from a through one column and collect the solution in a beaker. Repeat with the other ion exchange column using an equal volume of permanently hard water.

Class experiments
a Set up six test tubes in a rack, labelled A – F, containing about 1 cm depth of

     A Temporarily hard water
     B Permanently hard water
     C Temporarily hard water that has been boiled and filtered
     D Temporarily hard water that has passed through an ion exchange column
     E Permanently hard water that has passed through an ion exchange column
     F Distilled or deionised water

b Collect 10 cm³ of soap solution in a small beaker.

c Add a drop of soap solution to tube A. Stopper the tube and shake it. If no lather (foamy bubbles) appear, add another drop, stopper and shake again. Continue until a lather appears that lasts for 5 seconds or longer. Count the number of drops that you have used. Note the appearance of the water in the test-tube.

d Repeat the procedure in c for tubes B to F.

e In another test-tube, take a fresh sample of any of one of the water samples that were ‘hard’ (that is, those that took a lot of soap to achieve a lather). Add half a spatula measure of sodium carbonate crystals to the test-tube and shake it. Observe the contents of the test-tube. Now repeat the procedure in c to see how many drops of soap solution are required to produce a lather. Note the new number of drops.

f Repeat the procedure in e for any other water samples that were ‘hard’.

Teaching notes

A and B should require a lot of drops of soap solution, while the others should not require many at all. A and B contain dissolved calcium salts that react with soap solution to form an insoluble ‘scum’ that should be seen as a white cloudiness in the tubes or as specks floating on the surface of the water:

calcium salt(aq) + sodium stearate (soap)(aq) → calcium stearate (scum)(s) + sodium salt(aq)

Only when all the calcium ions have been precipitated out as scum will the water lather. Thus hard water wastes soap as well as causing unsightly deposits on baths and showers.

Temporarily hard water is defined as that which can be softened by boiling. The reactions by which it is made here are:

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

(Calcium carbonate is the ‘milkiness’ that forms when limewater is reacted with carbon dioxide)

CaCO₃(s) + CO₂(g) + H₂O(l) → Ca(HCO₃)₂(aq)

This reaction also occurs when rain water (containing dissolved carbon dioxide) flows over limestone rocks. On boiling, the reaction is reversed:

Ca(HCO₃)₂(aq) → CaCO₃(s) + CO₂(g) + H₂O(l)

The calcium carbonate shows as a white cloudiness (precipitate) when the temporarily hard water is boiled. The water does not now contain any dissolved calcium salts, so it is no longer hard.

This solid calcium carbonate is ‘limescale’ that wastes energy if it forms in boilers and kettles and can be dangerous if it blocks pipes or washing machines.

Hard water of both types can also be softened by:

• exchanging sodium ions for the calcium ions – these stay on the zeolite resin. This resin is a lattice with negative charges attached. These hold the positive ions. The attachment of the positive ions to the resin is reversible. The resin can be ‘regenerated’ by treating it with concentrated sodium chloride solution.
• adding sodium carbonate. This precipitates out the calcium ions as insoluble calcium carbonate:

eg CaSO₄(aq) + Na₂CO₃(aq) → CaCO₃(s) + Na₂SO₄(aq)

The calcium carbonate is once again seen as a white cloudiness.

‘Bath salts’ often contain sodium carbonate (as well as perfume etc) and this softens the water.

• ‘complexing’ the calcium ions – ie adding large anions that form a ‘complex’ with the calcium ions and stop them reacting with soap to form scum. Water softeners such as ‘Calgon’ work this way. The chemistry of complex ions is beyond the intermediate level, however.

Heating substances

Heating substances

Eye protection should be worn whenever a substance is heated.

Heating substances is always exciting but it is essential to keep the amounts used to a minimum and to use the correct apparatus in the recommended manner.

Although improvised containers can be used, the notes here refer only to the pieces of laboratory apparatus most commonly used for heating substances:

• ignition tubes (small test-tubes)
• test-tubes
• boiling tubes (large test-tubes)
• beakers
• evaporating basins
• crucibles.

Heating solids in test–tubes
Wear eye protection.

Only fill to a maximum of 1/5 ful                                 

Use a suitable test-tube holder
Hold the test-tube at a slight angle (see diagram)
Ensure that the open end of the test-tube isn’t pointing directly at anybody
Hold the test-tube so that the bottom is just in the tip of the flame
Always start heating with a small, gentle flame.
 

Heating liquids in test tubes
Wear eye protection.

Use a boiling tube (wide diameter) and do not fill to more than 1/10 ful        
Add an anti-bumping granule to give smoother boiling. Add the granulebefore starting to heat
Use a suitable holder
Hold the tube at an angle so that the top is well away from the flame
Hold the test-tube so that the bottom is just in the tip of the flame
Keep the liquid in the tube moving gently
For flammable liquids, use a water bath.

Heating flammable liquids

Wear eye protection                                                                                               

Use a boiling tube (wide diameter) and do not fill to more than 1/10 full.
Add an anti-bumping granule to give smoother boiling. Add the granule beforestarting to heat.
DO NOT heat directly over a naked flame.
Stand the tube in a beaker of hot water (e.g. from a kettle or hot tap).
 

Heating in beakers (and conical flasks)
Beakers should only be filled to 1/3 of their capacity when used for heating liquids. The addition of a few ‘anti-bumping’ granules will ensure smoother boiling.
 

Heating in evaporating basins
A flat-bottomed evaporating basin can be heated by supporting it on a wire gauze on a tripod. A round-bottomed evaporating basin is very unstable on a wire gauze so should be supported on a pipe-clay triangle when heating.

Evaporating basins should be filled to between 1/3 and 1/2 full.

When evaporating salt solutions, the solution should be heated (with occasional stirring using a glass rod), until solid just appears evenly around the edge of the liquid. The solution can then be left to cool - possibly overnight (labelled with names of the owners, the chemicals and any relevant safety warnings).
 

Heating in crucibles
A crucible must be heated on a pipe-clay triangle and not on a gauze. Start with a small, gentle flame before gradually increasing the heating rate. Allow plenty of time in the lesson for crucible and contents to cool down

Heating Group 1 metals in air and in chlorine

Heating Group 1 metals in air and in chlorine

Demonstration

This is a demonstration that shows the reactions of Group 1 metals in air and in chlorine. It does not clearly show the trends in reactivity of Group 1 metals, which is better demonstrated by the reactions in water, which follow on well from this demonstration.

To view a video of this demonstration experiments.
 

Lesson organisation

This experiment must be done as a demonstration. If you have not attempted this experiment before, it is strongly advised that you try it before performing the demonstration in front of students.

The first step is to generate chlorine, which can be done in advance, and requires a fume cupboard. The rest of the demonstration can be done in a well-ventilated laboratory. Goggles should be worn during the chlorine generation and the demonstration. The class should also wear eye protection during the demonstration.

How long the demonstration takes depends largely on how much talking you do between each part of the experiment. To speed things up a bit, the metals can be pre-cut into appropriately sized pieces, but they should be returned to the oil until just before they are used. For some classes it may be appropriate to do just one part of the experiment and heat the metals in either air or chlorine.

Chemicals

Lithium (HIGHLY FLAMMABLE, CORROSIVE) (Note 3 and 4)

Sodium (HIGHLY FLAMMABLE, CORROSIVE) (Note 3)

Potassium (HIGHLY FLAMMABLE, CORROSIVE) (Note 3)

Sodium chlorate(I) solution, 10-14% (w/v) (CORROSIVE), fresh (Note 2)

Hydrochloric acid, 5M (IRRITANT AT THIS DILUTION) (Note 2)

Apparatus

Goggles for the demonstrator, eye protection for the audience

Fume cupboard (only for generating the chlorine)

Clean, dry bricks with at least one flat surface, 3 (Note 1)

Gas jars with lids, 3 (Note 1)

Bunsen burner

Heat resistant mat

Scalpel

Forceps or tweezers

Tile

Filter paper

Universal Indicator paper 

Chlorine generator (TOXIC, DANGEROUS FOR THE ENVIRONMENT) See Standard Techniques: Generating collecting and testing gases (Note 2) 

Health and Safety and Technical notes

Demonstrator to wear goggles or face shield, class to wear eye protection. 

Chlorine - see CLEAPSS Hazcard. 

Lithium - see CLEAPSS Hazcard.

Sodium - see CLEAPSS Hazcard.

Potassium - see CLEAPSS Hazcard.

Sodium chlorate(I) - see CLEAPSS Hazcard. 

Hydrochloric acid - see CLEAPSS Hazcard and Recipe book.

1 The mouth of the gas jar must be narrower than the brick, to reduce the amount of gas escaping during the demionstration.

2 There are two methods given in the standard techniques for generating chlorine. The method that uses sodium chlorate(I) is safer than the method that uses potassium manganate(VII), but will not work well if the sodium chlorate(I) is an old sample as the concentration will be too low. Note that sodium chlorate(I), NaOCl, is NOT the same as chlorate(V), NaClO₃.

3 It is very helpful to have the 3 mm cubes of lithium, sodium and potassium cut ready to use. These should still be kept under oil until they are required.

4 Do not heat lithium in crucibles or other porcelain material – explosions have occurred.

Procedure

Heating the metals in air
a Starting with lithium, use the tweezers to pick up a small piece of metal and place it on a tile. Use a scalpel to cut a small cube with an edge of about 3 mm. Show the students the freshly cut surface which soon tarnishes, showing that the metal reacts quickly with oxygen. Blot off the oil using the filter paper, and place it onto the flat surface of the brick.

b Heat the metal from above using the hottest part of a roaring Bunsen flame just beyond the blue cone. Once the metal is on fire, remove the Bunsen flame. You should be able to observe the classic red of a lithium flame. (You may initially see a yellow flame, but this is the burning of any oil which was not removed.)

c Once the metal has stopped burning, test the residue with damp indicator paper and show that it is alkaline.

d Repeat for sodium and potassium.
 

Heating the metals in chlorine
a Prepare gas jars on chlorine in advance using a chlorine generator. See Standard Techniques: Generating collecting and testing gases.

b Check that the mouths of the gas jars of chlorine are narrower than the brick to reduce the amount of escaping gas, and that the colour of the gas in the jar is green. If it is not then there is not enough chlorine present for the demonstration to be successful.

c Starting with lithium, cut a small cube with an edge of about 3 mm. Blot off any excess oil. Place it on the clean, dry brick.

d Heat the piece of metal from above using the Bunsen burner (see diagram below). When the metal is burning, take away the Bunsen, invert the gas jar, remove the lid and immediately place over the burning metal. It helps to have a second pair of hands to do this. The metal continues to burn, producing fumes of white chloride. This method avoids producing FeCl₃ or CuCl₂, which can occur when a combustion spoon, or deflagration spoon made of iron or brass, is used.

e Repeat for sodium and potassium.

·         Heating sodium in chlorine Reproduced from CLEAPSS Guide: L195 'Safer Chemicals, Safer Reactions', section 9.3, by permission of CLEAPSS®

·         Heating sodium in chlorine Reproduced from CLEAPSS Guide: L195 'Safer Chemicals, Safer Reactions', section 9.3, by permission of CLEAPSS®

 Teaching notes

When heating in both air and chlorine, the expected pattern of lithium being the least reactive through to potassium being the most reactive may not be observed as it is hard to see potassium burning without the Bunsen flame. This may well be due to the potassium reacting faster than the others and an oxide coating being formed almost as soon as you begin to heat it.

In air:

4Li (s) + O₂ (g) → 2Li₂O (s)

Sodium and potassium produce a mixture of oxides, peroxides and superoxides.

In chlorine:

2Na (s) + Cl₂ (g) → 2NaCl (s) and similarly for the others.

The typical flame colours for lithium (red) and sodium (yellow) can usually be seen and sometimes the lilac of the potassium flame.